Sciencemadness Discussion Board

Hydrazine

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garage chemist - 28-4-2005 at 13:09

Sorry, starting from bleach just isn't worth it.
And my 5% old bleach contains way too much NaCl for strenghtening to be convenient. A lot of NaCl crystallizes out after addition of NaOH. The double amount would crystallize after strenghtening and turn the whole thing into slush.
And do you really think the bleach manufacturers use distilled water for their product? I don't.
My stirring rod was a bit brown on its side because I stirred a MnO2 suspension with it before. The yield was still quite high, as you've seen.
What was the purpose of the gelatin, after all?
I think the catalytic action of a hundreth ppm of metal ions is only true for the un- gelatinized reaction mix.
Metal ions still have to be avoided, but the reaction isn't hyper- sensitive.

To make my 10% NaOCl, I will have to chlorinate the NaOH until saturation, and then quickly add some extra NaOH (until the chlorine smell disappears) to stabilize the NaOCl. I hope this works.

To produce chlorine from TCCA, I simply add dilute (15- 20%) HCl. The reaction is smooth and steady (it can be controlled by warming or cooling, warming is usually necessary). Not only the chlorine of the TCCA, but also the chlorine of the HCl is converted to Cl2 gas, making this a very effective method.

Rosco Bodine - 28-4-2005 at 15:38

Quote:
Originally posted by garage chemist
Sorry, starting from bleach just isn't worth it.
And my 5% old bleach contains way too much NaCl for strenghtening to be convenient. A lot of NaCl crystallizes out after addition of NaOH. The double amount would crystallize after strenghtening and turn the whole thing into slush.


That is some really strange bleach . Are you speaking of the old 12% which has deteriorated ? Because that is the only way I can make sense of the untypical and disproportionate salt content .

Quote:

And do you really think the bleach manufacturers use distilled water for their product? I don't.


The water requirement is stringent for a stable commercial product . The half-life for regular bleach should be a year or more , and to get that requires highly pure water , if not distilled then exhaustively de-ionized . Often a second
flocculant like Mg(OH)2 is used , even done repeatedly if needed on the finished
bleach to remove any impurites which were introduced with the NaOH or from the equipment and holding tanks , as a finish treatment before bottling and shipment . Quality control is tight , or the product doesn't keep and comes back as defective .

Quote:

My stirring rod was a bit brown on its side because I stirred a MnO2 suspension with it before. The yield was still quite high, as you've seen.
What was the purpose of the gelatin, after all?
I think the catalytic action of a hundreth ppm of metal ions is only true for the un- gelatinized reaction mix.
Metal ions still have to be avoided, but the reaction isn't hyper- sensitive.
Quote:


The gelatine is used as a chelation agent to tie up metal ions , mainly copper is the worst for poisoning the hydrazine reaction . The metal ion data was in reference to the stability of the hypochlorite solution , where any impurity
is catalytic in the decomposition of the hypochlorite . But that won't be a problem for you since you won't be storing the solution . However it will
be a factor in the hydrazine synthesis
since what makes bleach unstable also
hinders the formation of the hydrazine .
Quote:

To make my 10% NaOCl, I will have to chlorinate the NaOH until saturation, and then quickly add some extra NaOH (until the chlorine smell disappears) to stabilize the NaOCl. I hope this works.


Keep it really cold towards the end of what you are planning to do . You are working contrary to the physical chemistry
which applies .
Quote:

To produce chlorine from TCCA, I simply add dilute (15- 20%) HCl. The reaction is smooth and steady (it can be controlled by warming or cooling, warming is usually necessary). Not only the chlorine of the TCCA, but also the chlorine of the HCl is converted to Cl2 gas, making this a very effective method.


That's an interesting method for chlorine .
TCCA has not received much attention in regards to its usefulness in synthesis .
That's a good one to know .

redneck - 29-4-2005 at 07:03

Hi guys
What recipe are you using? Where can I get it? Who is Mr A. , and where can I get his recipe?

@ Rosco Bodine
What is a mid file? I opened it with Winamp but I didn´t hear anything.

I don´t believe sodium hypochlorite is made the way you descibed on an industrial scale. I think it would be cheaper to make it via electrolysis of an cold aquaeous solution of sodium chloride.

Rosco Bodine - 29-4-2005 at 07:36

Quote:
Originally posted by redneck
Hi guys
What recipe are you using? Where can I get it?


http://www.sciencemadness.org/talk/viewthread.php?tid=757&*
Quote:

Who is Mr A. , and where can I get his recipe?


Speak of the devil , ........
and then look above :D
Quote:

@ Rosco Bodine
What is a mid file? I opened it with Winamp but I didn´t hear anything.


Try Winblows media player and enable midi as a default file type association .
Quote:

I don´t believe sodium hypochlorite is made the way you descibed on an industrial scale. I think it would be cheaper to make it via electrolysis of an cold aquaeous solution of sodium chloride.


Actually the same plant usually produces three products , sodium hydroxide , chlorine , and sodium hypochlorite .
It allows for continuous operation , and diverting production into what is sold separately or already combined .

redneck - 29-4-2005 at 09:45

@ Rosco Bodine
Thank you for the link!

Strange music. Sounds like south american indian style. I prefer Carl Cox when I´m working. He helps against tiredness.


I tried to make hydrazine sulphate from ammonia. I read recipes in the internet which use 2.5 to 25 times more ammonia than stochiometrically needed.
But I didn´t want to have a terrible smell smell in the garage again. So I used not more ammonia than theoretically needed.

@ garage chemist
The chlorox I have is labelled as containing 2.8 % NaOCl. But when I boiled it down untill dryness I received 13 % of an unknown salt. If it was NaOCl*6 H2O the chlorox would contain 5.7 % NaOCl.
I also tested how much chlorine gas is generated when I add HCl 30 % to the chlorox. From the amount of gas I calculated the amount of NaOCl in the chlorox. My result was 6 %.
So I figured out that the label was wrong. The chlorox contains not 2.8 % NaOCl but 2.8% Cl.
So the chlorox contains 5.88 % NaOCl for real.



My hydrazine sulphate experiment

2 NH3 + NaOCl ==> N2H4 + NaCl + H2O

I poured 20 ml chlorox into a test tube. Then I dissolved 13 mg gelatine and 36 mg sodium acetate (It´s also a heavy metall ion catcher) in it. Then I cooled it in a water bath and slowly added 2.2 ml 25 % ammonia. The tube would become warm by the reaction heat and the ammonia would evaporate without cooling. During the addition of the ammonia there was some foaming and after the addition a tiny but steady stream of bubbles coming out of the mixture. Then I added 36 mg dirty urea and boiled the solution down to one third of it´s original volume.
Then I let it cool and afterwards I cooled it in a water bath and added 3.2 ml 38 % sulfuric acid in small portions because a lot of gas was generated and it foamed. When half of the acid was added there were tiny crystals in the solution. I filtered them of and added the rest of the acid to the solution but no more crystals were formed. But the next day there were large crystals in the size of one inch long needles in the solution. But there were also medium size cubic crystalls in the solution which look like table salt.

Now the question is are these needles hydrazine sulfate?
Are the cubic crystalls NaCl or are they also hydrazine sulfate?
How can I find out what they are?
What are the tiny crystals that formed immediatly when I added the sulfuric acid?Can I use more gelatine without turning the solution into a pudding like mass?
(How many per cent compared to the boiled down solution)

Could anybody post how hydrazine sulphate crystals look, please. What is their shape?

[Edited on 29-4-2005 by redneck]

12AX7 - 29-4-2005 at 11:40

Boiling NaOCl breaks it down into ClO3 + 'Cl.

Doesn't adding ammonia to bleach release chlorine instead? Or am I thinking of acid?

Tim

redneck - 29-4-2005 at 11:55

3 NaOCl ==H+, >70°C==> NaClo3 + 2 NaCl

NaOCl converts to NaClO3 only if acidified. In an alkaline solution it´s stabe.

NaOCl forms chlorine gas when it reacts with HCl not with ammonia.

sparkgap - 29-4-2005 at 19:36

Ammonia and bleach give chloramine (Cl-NH<sub>2</sub>;), Tim.

Rosco, why don't you just tell everybody else who Mr. Anonymous is? ;) :P

Somewhat related to the topic: how long will stored hydrazoic acid last before decomposition?

sparky (~_~)

Rosco Bodine - 3-5-2005 at 19:19

Quote:
Originally posted by sparkgap

Rosco, why don't you just tell everybody else who Mr. Anonymous is? ;) :P


Mr. Anonymous and Rosco Bodine are aliases . These aliases conceal
the true identity of Rosco P. Coaltrain , sheriff of Hazzard County ,
to keep him out of trouble with Boss Hogg , chief of animal farmland security :D
Quote:

Somewhat related to the topic: how long will stored hydrazoic acid last before decomposition?

sparky (~_~)


The dilute solution is reportedly stable in the absence of light and kept cool .
But such solutions are an extreme danger due to the toxicity and volatility .
The material is so toxic that it has been said if you smell it , then you have
already been poisoned . And that is not very much if any exaggeration .
Extreme Caution is the imperative rule for hydrazoic acid . It simply is not
the sort of material inclined to suffer fools or carelessness without exacting
a dear price for the error . It is among the deadliest toxins known to science .

sparkgap - 3-5-2005 at 20:41

The container of hydrazoic acid I saw in our lab was being kept at quite cool temperatures, so I guess its volatility and rate of decomposition should be reduced. Still, I'm thankful I've never needed to use it in any of my experiments.

I thought as much about the toxicity. :) IIRC, azide ion is a good inhibitor of the ETC so crucially needed in respiration.

Thanks, Mr. Anonymous. :D

sparky (^_^)

Rosco Bodine - 3-5-2005 at 21:14

The recent discussion about chlorination of NaOH has caused me to remember that this came up once before , and I will have to check my patents , but IIRC there is a more efficient method for use of chlorination in the synthesis of hydrazine and it involves chlorination of a slurry of urea to form chlorourea , which is then
reacted with a base like sodium carbonate
or sodium hydroxide to form hydrazine upon heating and hydrolysis . IIRC , there is no need to even make sodium hypochlorite , if you are going to go the chlorination route , and the yield is better by the chlorourea intermediate . I will see if I can find the patents which reference the method .

GB1063893 Hydrazine or Semicarbazide selectively produced from monochlorourea

Related semicarbazide patents are
US5241117
GB1153483

[Edited on 4-5-2005 by Rosco Bodine]

garage chemist - 4-5-2005 at 06:11

Wow, that chlorourea route would be great if it worked! Thanks for the suggestion.
It reminds me of a method to produce methylamine: first acetamide is mixed with bromine, then the formed N-bromoacetamide is treated with a base to form methylamine.
That should work with urea, too!
I have about 7ml of bromine, so I could try it out.
A method for determining if a precipitate which occurs on adding H2SO4 is indeed HS, one could dissolve it in some ammoniacal AgNO3 solution. HS immediately produces a silver mirror, without any warming.

[Edited on 4-5-2005 by garage chemist]

Rosco Bodine - 4-5-2005 at 07:00

Earlier in the thread a couple of mentions were made of the urea - hypochlorite method having been performed successfully using lower concentration bleach . I have never tried the use of the lower concentration simply because I was seeking to obtain the greatest amount of product per liter of reaction mixture which would be processed . But there is a German patent which refers to what is probably the invention of this
" Schestakow reaction " for hydrazine from urea and hypochlorite , and the original patent describes dilute hypochlorite of 3% being used , if I have understood the German text correctly .
See attached file .

Anyway , the conditions under which hydrazine forms and can be isolated from such reactions are very specific . In different molar proportions and conditions of pH and temperature and reaction time ,
indeed urea and hypochlorite mutually decompose each other without the formation of hydrazine . And this fact makes urea a patented decontaminant used for the environmentally safe destruction of waste hypochlorite solutions . So from this it is known that the manufacture of hydrazine is a very specific and controlled process for a narrow range of conditions . Departure
from that " window " reaction condition results in yields of hydrazine rapidly deteriorating to zero . This fact is why I
tried to emphasize earlier that changing anything about the " optimized " method would reduce the yields . For information
related to the conditions for urea being used for the destruction of hypochlorite see US4508697 , US2784057 .

An off topic digression here concerning urea and a possible reaction product with sulfur , hydrogen sulfide , or sodium hydrosulfide / sodium polysulfide . I cannot recall for certain where I read a passing reference concerning such a reaction , or even if it is perhaps just something in the way of a contemplated experiment which occurred to me out of the blue . It was an idea in regards to producing a fusible and or soluble organo-sulfur compound which may have usefulness as a fuel in propellants , or as a fuel/sensitizer for AN compositions , or as a possible sulfide type reducing agent .
I have searched exhaustively and can find no reference to such a urea reaction . If anyone knows of something like this or
a reference , please refresh my memory .

[Edited on 4-5-2005 by Rosco Bodine]

Attachment: DE164755 Schestakow , Hydrazine by Urea Hypochlorite.pdf (157kB)
This file has been downloaded 2164 times

garage chemist - 4-5-2005 at 15:08

Thanks a ton for the german patent. Too bad I can't read it at the moment... my computer crashes as soon as it has fully loaded. That's my old computer... my better one is being repaired at the moment.

Anyway, today I have sucessfully produced sodium azide.
I did the freebasing with ethanol since I was running low on IPA and ethanol is just easier to get.
To 13g HS, I added 18ml denatured ethanol (94% Vol.). Then 4g of NaOH were added.
Upon stirring with a glass rod, the salt mixture became dough- like and eventually turned into a homogenous sticky white mass. The mixture was quite hot. I wanted to proceed quickly in order to not lose too much hydrazine due to air oxidation.
I added another 5g NaOH, and suddenly the mixture started boiling violently. White fumes were produced, and some of the solution was spilled. I quickly added a dash of ethanol to cool it. This stopped the boiling for a moment, but it started boiling again as I kneaded the mass with the stirring rod.
I had to add even more ethanol.
But finally, only a dry, very fine and rather voluminous precipitate of sodium sulfate remained under the ethanol.
The solution was decanted and the sodium sulfate washed three times with 10ml ethanol each time (the precipitate sucked up a lot of ethanol and was rather voluminous).
The combined solutions were stored in a tightly closed round bottom flask.

5,5g NaOH were dissolved in 60ml ethanol (took quite a lot of time). This was then mixed with the hydrazine extract. It was nearly 100ml of liquid.

18,7g IPN were prepared from 15g NaNO2, 17g IPA and 40ml 25% HCl (an excess of IPA was used in order to obtain a stable IPN). The NaNO2 was dissolved in a minimal amount of water and the IPA was added. Then, while cooling in an ice bath, the HCL was slowly added. The IPN yield was very good (over 90%). It was washed with bicarbonate solution and "dried" for 15 minutes with CaCl2 (I could have omitted that step, it was not necessary).

The ethanolic NaOH/hydrazine solution was placed in an ice bath and one third of the IPN was added. After 15minutes of stirring, there was no evidence of any reaction. The second third was added. Again no reaction. The last third was added, again without any signs of precipitate forming. I took the solution out of the ice bath and waited for it to warm up. Finally, after 30min, it became more cloudy and an extremely fine precipitate began forming. It became a bit warm and I stuck it into the ice bath again.
The amount of precipitate slowly increased. I am letting it stand overnight in order to complete the reaction. The precipitate looks "milky", it is really very fine.

I observed the following things:

- The freebasing is VERY exothermic! After adding the first portion of the NaOH, the solution MUST return to room temperature before adding the second portion. The container should be covered during the cooling.
And 18ml of ethanol is not enough. 25ml would more reasonable.

- The diazotation reaction took an awful lot of time to start and then only proceeded lazily. I think this was due to the rather dilute solutions I used. I should have used less ethanol to dissolve the NaOH. And I should give the sodium sulfate from the freebasing more time to settle so I can use less ethanol for the washings.

That's how one learns! My next azide batch will be much better and without incidents like the hydrazine extract starting to boil.

[Edited on 4-5-2005 by garage chemist]

Rosco Bodine - 4-5-2005 at 15:25

When I do the freebasing which will later be extracted with methanol , the process
of reacting the solids first to form an aqueous phase is done before any alcohol is added . This allows for the exotherm to be managed , and actually it is essential to maintain
that exotherm at nearly the boiling point to keep the slurry liquid . The slurry is only allowed to cool just to the boiling point of the alcohol so it won't flash boil the alcohol when it is added . The first alcohol extract is still very warm when it is decanted . This works fine for methanol but not for isopropanol . I am unsure about ethanol
but I seem to recall that it worked okay and didn't phase separate like isopropanol .
Read my earlier post about the alternating additions which really helps the solubility for the freebasing .

You really need to be doing the freebasing in an Erlenmeyer with a stirbar on a magnetic stirrer . A glass stoppered
flask is best . The NaOH should be preweighed and quickly placed into a small plastic bottle having a capped conical dispenser tip , like a mustard bottle . The stirbar should be set at slow speed and about half the HS put into the flask . To the stirred dry mixture can be added in portions the NaOH , stoppering the flask and capping the NaOH bottle between additions . If you follow this method , you should get good results .
A drop or two of water may be required to kickstart the reaction which then produces its own water as it proceeds ,
melting the solids in the byproduct water
which self heats as it goes . What you are actually getting as the liquid , is hydrazine dihydrate , plus any water you may have to add , the less the better of course .

[Edited on 4-5-2005 by Rosco Bodine]

Just a writeup

BromicAcid - 5-5-2005 at 19:49

12 grams of sodium hydroxide was dissolved in 35 ml distilled water and added with the aid of a sepretory funnel (tight fitting ground glass) with magnetic stirring to 20 grams of hydrazine sulfate. The additions were done over the course of and hour and a half during which very slight heating was observed.

At no point did everything become liquid but I assumed the heavy precipitate at the bottom was sodium sulfate that had been produced. This mixture as then set up to distill. The apparatus is in the attached picture. I distilled into a sepretory funnel so that I could take off the distillate at different times. The gasses were simply lead away and bubbled into water.

The distillate started to come over when the water bath was at about 130 C and the still head only rose to 100 C so it was mostly water coming over. The bath temperature continued to climb until it held at 152 C and the stillhead continued to hold at 100 C. A portion of the liquid distilling was removed and a few drops were added to a test tube of bromine water, immediately removing the color. So a positive test for reducing character. A large portion of the hydrazine came over in this region, too much water, too afraid of distilling anything less then the azeotrope (hydrazine hydrate) at 120 C. It took a lone time to remove the water, the amount of azeotrope expected to come over was only 6.75 ml and therefore it came as no suprise when after 17 ml of water came over I only had a short break around the point of the azeotrope before there was a temperature drop and liquid stopped distiling. Additionally the flask was bumping quite a bit and I didn't want to press it considering it was hydrazine.

I know there are better ways to get the hydrate but this way has always bekoned to me, now I only have 480 g hydrazine sulfate left....

BTW, does anyone have any details on the preparation of anhydrous hydrazine by distilling it from hydrazine borate? Expecially the preparation of said compound?

[Edited on 5/6/2005 by BromicAcid]

blurrydistill.jpg - 65kB

Marvin - 6-5-2005 at 09:17

I have a reference but only the obvious directions.

"Reaction of the hydrate with boric acid gives hydrazinium borate which can be dehydrated by heat and decomposed to N2H4 and B2O3 by furthur heating."

And the reference quoted with it is Gmelins Handbuch der anorganische Chemie, Berlin, Verlag Chemie, 1928. 8th ed., System 21: p313-314.

I'd be interested to read that myself.

Rosco Bodine - 6-5-2005 at 10:41

Cyanuric acid similarly forms a salt with hydrazine which decomposes at high temperature to the free cyanuric aicd and free hydrazine . But the decomposition temperature is not stated , except as being below 360 C , which is damned hot .

Any of the distillation schemes for anhydrous hydrazine , even for the hydrate actually , are going to result in decomposition losses . The ketazines may be thermally unstable , and this is a possible route . I know the ketazines are
hydrolyzed to the salt by acids , and perhaps they may also be hydrolyzed by
gaseous ammonia , which could be another possible route for the hydrazine monohydrate . The hydrate can be dehydrated by highly activated and dry
alumina , but the adsorption is very small before the alumina has to be high fired again to get it dry enough for recycling .
It's only good for something like 3% of the weight of the alumina in absorbed water , till the deydrating property is lost .
Trying to get the water out of hydrazine
is something like trying to get the water out of sulfuric or phosphoric acid . It is
extremely difficult , and that is precisely why it is best to work with the freebasing method . Obtaining the hydrazine in a highly concentrated and dehydrated form
is a process plagued by losses and difficulty which should be avoided for any scenario where the hydrazine is not required for synthesis to be used in such concentrated form .

Marvin - 6-5-2005 at 11:58

I think anyone desperate enough to try that with ketazines would be better off doing ammonolysis directly on hydrazine sulphate with liquid ammonia or by using an entraining method on hydrazine hydrate. I agree with you though, anhydrous hydrazine is seldom useful and rather hard won.

garage chemist - 6-5-2005 at 14:30

I was able to open the PDF by first saving it and then opening from there.

The Schestakow process actually uses 3% NaOCl on urea, but the hydrazine is isolated by addition of benzaldehyde to the reaction mixture, benzaldazine is formed and this is seperated and hydrolysed to benzaldehyde and hydrazine sulfate with sulfuric acid.

I am unsure if the solubility of hydrazine sulfate is low enough to allow it to crystallize out from the dilute solution obtained by using 3% NaOCl.
Does anyone know the solubility of HS at 0°C? The solubility is always listed for 25°C.

EDIT: I will also try out the chlorourea process.

[Edited on 6-5-2005 by garage chemist]

S.C. Wack - 6-5-2005 at 15:01

System 21 of Gmelin's is sodium. The original article that Gmelin's is supposed to point to may be J. Russ. Phys. Chem. Soc. 34, 227 (1902), by Djavachoff, and the temp is over 260C to form hydrazine in that article.

As for Roscoe's complaint in the azide thread about the move (of discussion of preparation of hydrazine from hypochlorite and urea or whatever) from the azide thread to this one - I should comment since I'm the one who wondered how many more pages of azide thread the preparation of hydrazine was going to take. It seemed obvious to me that two full pages of this with no end in sight (see DPPP thread) was out of place, given that there was already this thread, and it would certainly be of more use here.

Rosco Bodine - 6-5-2005 at 15:50

Quote:
Originally posted by S.C. Wack

As for Roscoe's complaint in the azide thread about the move (of discussion of preparation of hydrazine from hypochlorite and urea or whatever) from the azide thread to this one - I should comment since I'm the one who wondered how many more pages of azide thread the preparation of hydrazine was going to take. It seemed obvious to me that two full pages of this with no end in sight (see DPPP thread) was out of place, given that there was already this thread, and it would certainly be of more use here.


Actually , the matter of the off topic nature of the preparation of hydrazine sulfate
had already been noticed , and mentioned before you posted the question
about how many more hydrazine posts in an azide thread would there be . And I fully agree that the preparation of hydrazine sulfate is off topic in the azide thread . Splitting the thread and exporting the synthesis of hydrazine sulfate posts here was okay by me .
No problem there .

But a half dozen posts that were azide specific were mixed in with the posts which were exported here from the other thread . What I meant to say is it was more than just the offending
off topic posts which got moved . I shouldn't have said anything because I bet I just volunteered to be the one to sort through those posts and pick out the candidates for
" reintegration " .

chemoleo - 6-5-2005 at 19:51

Hmm, it's a tricky one since things are so interlinked - the production of Hydrazine and azide - particularly when a single post contains descriptions of both (something I cant split easily without disrupting the flow).
If there's a specific post that I overlooked and doesn't belong here, let me know.

Suggestions welcome. I would only want to perfect the layout of such a high-level discussion :)

I will remove this post later, once/if I got suggestions (U2U pls).

PS oh yeah rosco you are damn right :D

[Edited on 7-5-2005 by chemoleo]

garage chemist - 11-5-2005 at 07:19

I did not have good sucess with the alcohol- free freebasing.
I mixed 13 HS with 4g NaOH, it got warm and after 2 hours with occasional stirring the mix slowly liquefied.
But on addition of the next 5g NaOH, nothing happened, the mix cooled down and the NaOH showed absolutely no signs of reaction.
I extracted it with ethanol three times and there were still unreacted portions (clumps) of the powder mix, even though I stirred it a lot.

With the addition of alcohol to the HS at the beginning (my previous method), I got a complete reaction and a fine precipitate of sodium sulfate.
If it can be cooled efficiently, this method would be optimal.

[Edited on 11-5-2005 by garage chemist]

Rosco Bodine - 11-5-2005 at 09:07

That's strange . I suspect your hydrazine sulfate is not pure ,
maybe a lot of salt or sodium sulfate as an impurity .

The freebasing generally is a very energetic and exothermic process and
proceeds rapidly . Five minutes , not two hours for a tenth mole of HS
would be more like it . And the exotherm proceeds
right to the end or almost the end . You really have little to no free hydrazine appearing until half of the NaOH has been added , since all you are doing is converting the hydrazine sulfate to dihydrazine sulfate to that point , if there is good mixing . So you can add the first half of the NaOH as rapidly as the exotherm allows and in fact the idea
is to get the mixture very hot , so everything dissolves as a hot concentrated solution of dihydrazine sulfate . Only during the additions of the remaining NaOH does the actual hydrazine freebasing occur , simultaneously with the precipitation of
anhydrous sodium sulfate forming a slurry
in a hot liquid phase which is decreasingly
dihydrazine sulfate , and increasingly converted to hydrazine dihydrate
by the time the final amount of NaOH is added . IIRC the freebasing of 2 moles of HS , took perhaps a half hour to the point where the first of the alcohol ( methanol )
was added . And a smaller mass like a
tenth mole would likely require supplemental heating , not cooling ,
in order to maintain everything in solution
at the midpoint of the reaction , before the actual freebasing part of the reaction begins with the addition of the remaining NaOH . You can follow the course of the reaction by observing the the exotherm and the solubility changes which are very predictable .

garage chemist - 11-5-2005 at 12:43

The problem may have been that my NaOH was in the form of rather large granules (2-3mm) and also the very small amount, which I should have heated instead of cooled.

My HS is not impure homever, as the freebasing in ethanol is extremely exothermic, as I've described earlier.

Freebasing with ethanol is just a better method for small batches, and that's what I want.

Larger amounts might work better without ethanol, but I'm working on a small scale here.

BTW, about azides again: for 0,1 mol Hydrazine, I use 0,18- 0,2 mol IPN. Is that okay? (No nitrite is lost via transesterification, as I keep the reaction vessel tightly closed)

Today I used the hydrazine extract to make another batch of NaN3, this time with only 50ml of reaction mixture, and the reaction set in much more rapidly. I added one third of the IPN to the ice- cold hydrazine/NaOH solution and took it out of the ice bath. It became cloudy after about 15 minutes of warming up. I've put it into the ice bath again and added the whole rest of IPN at once (all with strong stirring). Then I tightly closed the flask and I'm leaving it in the ice bath over night. The ice bath slowly warms up and allows the reaction to proceed slowly.

[Edited on 11-5-2005 by garage chemist]

Rosco Bodine - 11-5-2005 at 14:25

Quote:
Originally posted by garage chemist
The problem may have been that my NaOH was in the form of rather large granules (2-3mm) and also the very small amount, which I should have heated instead of cooled.


Yes that is the reason for the sluggish reaction , the state of subdivision of the
reactant is limiting the contact area .
With that coarse material , it may hurt nothing at all to simply dump it all in together at once , even to kick it a little with a few drops only of water .

The granulation of the NaOH which I have been using is very fine , a little more coarse than sugar . The fine granulation
reacts on contact and completes in seconds . So your slow reaction is from the rock salt like chunks being slow to dissolve .
Quote:

My HS is not impure homever, as the freebasing in ethanol is extremely exothermic, as I've described earlier.

Freebasing with ethanol is just a better method for small batches, and that's what I want.


As a general conclusion that is premature . I still think that at least to the point of the dihydrazine sulfate intermediate the aqueous system would prove most efficient . A little heat may change the result you are getting . But
if the yields are good the way you are doing it then that is what counts .
Quote:

Larger amounts might work better without ethanol, but I'm working on a small scale here.


Yes , the scale and concentration and
what you are trying to do within a given volume influences your choice of method ,
and the thermal considerations are no small factor . I tend to work with what
some would consider huge batches , but
having an industrial background has distorted my perspective towards molar
instead of millimolar quantities , unless caution about some stabilty issue or toxicity finds me working with "mini" scale . I dislike working with very small quantities when doing reactions about which I have confidence , seeing the same amount of labor and steps required can be scaled up greatly and then the " stash " of the desired material
is increased , having made the effort more
worthwhile . To me a modest batch of NaN3 is 75 grams or more from a liter of
methanolic Hydrazine Extract :D

Quote:

BTW, about azides again: for 0,1 mol Hydrazine, I use 0,18- 0,2 mol IPN. Is that okay? (No nitrite is lost via transesterification, as I keep the reaction vessel tightly closed)


Actually you don't want too much excess of nitrite , as it will decompose the azide already produced and reduce the yields .
I wouldn't go beyond 15-25% above theoretical on the nitrite if using IPN ,
or ethylene glycol dinitrite , for nitrosation
of an ethanolic extract . And you don't want the reaction mixture too cold , nor would it be recommended to have a sealed container with no overpressure relief , because IIRC you can get nitrogen buildup from this reaction .
Quote:

Today I used the hydrazine extract to make another batch of NaN3, this time with only 50ml of reaction mixture, and the reaction set in much more rapidly. I added one third of the IPN to the ice- cold hydrazine/NaOH solution and took it out of the ice bath. It became cloudy after about 15 minutes of warming up. I've put it into the ice bath again and added the whole rest of IPN at once (all with strong stirring). Then I tightly closed the flask and I'm leaving it in the ice bath over night. The ice bath slowly warms up and allows the reaction to proceed slowly.

[Edited on 11-5-2005 by garage chemist]


What I have observed with the reaction run in methanol is that there definitely is an induction period for the reaction before
the sodium azide begins to precipitate ,
and it may be that the water present causes the initial product to be held in solution until saturated and then precipitation begins , and continues at about the same rate as nitrite is added .
The crystals will be a bit more coarse I think with gradual addition of the nitrite ,
to the stirred mixture . If you save the residual alcohol solution from filtration ,
you can bring it to boiling to drive off
some of the water adding some makeup
alcohol to the mixture as needed , and after cooling you can add maybe one fourth the volume of toluene , and a second crop of less pure azide will precipitate from the cloudy solution .

garage chemist - 12-5-2005 at 13:31

ARRGH!!! I only washed the NaN3 with denatured ethanol (non- purified, since I thought that all the hydrazine had reacted) and the ketazine formed again! An oily blob is in the filtrate and the NaN3 leaves a black residue after deflagration!

This is a major drawback of the diazotation in ethanol.
The (purified!) ethanol also turned yellow
during the diazotation, indicating that it was not pure.

The strange thing is that the hydrazine extract with NaOH did not form any precipitate and stayed totally clear before the IPN was added (I left it to stand overnight)!
The Nitrite also played a role, I think.
Or the azide?
I'm sure that there is some other junk in my "ethanol" besides the standard MEK denaturant.

I really should switch brands or better buy some absolute ethanol denatured with petrol ether or toluene.

PS: Thanks a lot for your very informative posts, Rosco! Your help is much appreciated.

[Edited on 12-5-2005 by garage chemist]

Rosco Bodine - 12-5-2005 at 15:52

Quote:
Originally posted by garage chemist

The (purified!) ethanol also turned yellow
during the diazotation, indicating that it was not pure.
Quote:


The yellow color to the reaction mixture is normal when organic nitrite is present in the cold mixture . After you estimate the reaction in the cold is completed , you can heat the mixture and drive off the nitrite in solution , and the mixture will clarify .

Quote:

The strange thing is that the hydrazine extract with NaOH did not form any precipitate and stayed totally clear before the IPN was added (I left it to stand overnight)!


Probably there is little carryover of any carbonate or sulfate when the alcohol extraction is not done hot . Or else these
impurities are less soluble in ethanol .
Try storing the extract in the freezer overnight and see if any small precipitate
settles out . It is for me a very small amount , which settles onto the glass .
Quote:

The Nitrite also played a role, I think.
Or the azide?
I'm sure that there is some other junk in my "ethanol" besides the standard MEK denaturant.


The most common denaturant is methanol . No problem with any azine formation is there . Though you will lose some nitrite due to formation of methyl nitrite , it shouldn't be so much a problem since the concentration of methanol is low . In a cold mixture a lot of the methyl nitrite will actually react before it is lost .
Quote:

I really should switch brands or better buy some absolute ethanol denatured with petrol ether or toluene.


Check the label or the MSDS for the product and it should say what denaturant is used . Honestly I didn't think of the azine formation when I suggested a non-methanol denatured ethanol . Sorry for that oversight . The methanol denaturant would be preferable
to a ketone denaturant . My foulup there .

S.C. Wack - 12-5-2005 at 19:17

Found the Gmelin's reference mentioned earlier in 23, 565 (1936):

Hydrazoniumborat (N2H5)2H4(B4O7)3. Beim Neutralisieren einer Lsg. von krystallisierter
Borsäure mit 50%iger N2H5OH-Lsg., Einengen auf dem Wasserbad und über H2S04 erhält man rhomb.
Krystalle des 10-Hydrats (N2H5)2H4(B4O7)3.10H2O. Verwittert leicht und geht im Vak. über H2SO4
in das 5-Hydrat (N2H5)2H4(B4O7)3.5 H2O über. Beim Erhitzen auf 100° entsteht das wasserfreie
Salz, bei 250° bis 260° (N2H4)2(B2O3)6; oberhalb 260° schmilzt die Subst. und zersetzt sich dabei in
B203 und N2H4, A. Dshawachow (J. Russ. Ges. [chem.] 34 [1902] 227; C. 1902 I 1393).

"C." stands for Chemisches Central-blatt, perhaps that abstract has more detail, short of someone finding the Russian article.

garage chemist - 17-5-2005 at 11:57

Thanks, S.C. Wack. This seems to be the easiest way to make anhydrous hydrazine.

Today, I tried to make HS with 2,8% NaOCl.
The amounts of NaOH, urea, gelatin, HCl and H2SO4 were adjusted according to the NaOCl content.
On adding the urea/gelatin to the basified NaOCl at 10°C, nothing happened. No foam, no heating. I slowly heated the solution to 60°C and held it there for 30 minutes (the color was a bit yellowish), then I heated it to the boling point (it turned almost colorless).
I cooled it down and added the calculated amount of HCl and the calculated amount of H2SO4 (only very little gas was evolved). Nothing precipitated.
I checked the pH of the solution: VERY alkaline. Strange! I added more HCl until it was acidic. LOTS of gas and foam were now produced!
On cooling, still nothing precipitated. There is a very small amount (about 0,1g) fluffy "precipitate" which doesn't settle.

This was a failure.
My next experiment will be the production of 10% NaOCl from NaOH and chlorine and then production of HS from this.

EDIT: and if this doesn't work either, I'll try the chlorourea process.

[Edited on 17-5-2005 by garage chemist]

Rosco Bodine - 17-5-2005 at 18:45

If chlorination is going to be a step in your
hydrazine synthesis , you should try the chlorourea method first ,
since it simplifies the rest of the synthesis and should produce a higher yield than the alternative . The hypochlorite and urea method is already well known to work , so experimenting with the relatively little known chlorourea method would make a more interesting experiment . This has the chance of confirming a method
barely mentioned in the literature , while
chlorination of NaOH is routine .

garage chemist - 18-5-2005 at 03:19

OK, all right.

Could there be some unknown hazards with the chlorourea process? The worst- case scenario would be the formation of nitrogen trichloride.:o
Interestingly, NaOCl + urea also yields large amounts of NCl3, but only under special conditions (NaOCl added to a mixture of urea and sulfuric acid- here, chlorine is first produced from the NaOCl and acid). There was a thread about this somewhere...

I'll try the chlorination of urea with very small amounts first and look if it's safe.

Mongo Blongo - 18-5-2005 at 07:41

Garage Chemist-I find that I have to heat the reaction with urea/hypochlorite from begining to end or the reaction will fail. This may have something to do with the low conc of the sodium hypochlorite used (it just says "less than 5%" on the bottle).
On another note, can anyone see a reason why acetone would be a bad choice to extract the hydrazine? I'm just thinking that it's easier to find pure than ethanol/methanol and probably won't give a phase separation like isopropyl will.

garage chemist - 18-5-2005 at 07:44

Acetone certainly won't work, because it will react with the hydrazine. Actually, my problem were the trace amounts of ketones in my ethanol!

froot - 18-5-2005 at 12:10

Another factor that could be attacking your yields of HS is when using low concentrations of NaOCL there is too much water in the mixture. HS is to some extent soluble in water and will remain in solution when the reaction is complete. This is where it is also imperative to follow a procedure such as Mr A's to the T. I may have overlooked this being mentioned earlier, if so, oops! I have used Mr A's procedure and it worked every time.

froot - 18-5-2005 at 12:12

Another factor that could be attacking your yields of HS is when using low concentrations of NaOCL there is too much water in the mixture. HS is to some extent soluble in water and will remain in solution when the reaction is complete. This is where it is also imperative to follow a procedure such as Mr A's to the T. I may have overlooked this being mentioned earlier, if so, oops! I have used Mr A's procedure and it worked every time.

garage chemist - 18-5-2005 at 14:58

@ Mongo Blongo: Do you use DanKlorix? It is 2,8% NaOCl, it is written on the bottle (under "Anwendungen", it's rather hidden).
I was not able to get any HS with 2,8% NaOCl.

Anyway, I'll never use hypochlorite again, because today I tried the chlorourea process and IT WORKED!!!

Here's what I did:

In a 3-neck 250ml rbf equipped with a gas inlet tube whose tip had been pulled out into a "pipette tip" (the other two necks were simply left open), I put 20,5g urea.
In a beaker, I dissolved 0,2g gelatin (this is still necessary!) in 45ml warm distilled water.
This solution was added to the urea and it was stirred until the urea had dissolved.
A gas generator was set up (100ml rbf and pressure- equalized dropping funnel with gas outlet at the side) and charged with 10g TCCA (made into a slurry with 10ml water) and 34ml 25% HCl.
The HCl was slowly added to the TCCA, the chlorine bubbled through the urea/gelatin. The urea solution and TCCA were swirled from time to time to ensure complete reaction. The chlorine bubbles reacted completely, no chlorine smell was noticed.
When all the HCl was added, the TCCA/HCl sludge was heated to nearly the boiling point to expel all the residual chlorine.

In a 500ml rbf, 16g NaOH were dissolved in 40ml distilled water. To the resulting hot solution the chlorourea solution was added immediately and all at once. It was swirled to ensure complete mixing.
The solution went dark orange and white foam was produced, the solution approximately doubled its volume.
It looked EXACTLY like a normal batch of HS with hypochlorite!
It was slowly heated, the foam slowly subsided and the color became lighter.
When it was almost boiling, it was nearly colorless. It was left to cool.

While stirring with a magnetic stirrer, 60ml 25% HCl (a bit more than needed to neutralise the used NaOH) were slowly dripped in, CO2 was being produced, and the hydrazine fumes+HCl gas formed a white fog in the rbf.
18ml of conc. H2SO4 were diluted with 20ml water and slowly dripped into the well stirred mix. After about two thirds were added, a precipitate suddenly occured.
The rest of the H2SO4 was added more slowly in order to produce larger crystals.

It was left to cool to room temperature, then it was cooled to 0°C in an ice bath. 100ml of the supernanant liquid were set aside to wash residual crystals out of the rbf.
It was filtered through a coffee filter and washed with a small amount of ice- cold distilled water and then with some ethanol.

It's drying at the moment. The theoretical yield (100%) would be 13g, but some HS stays in solution of course.
Tomorrow I'll weigh it. It looks like about 8 grams.

I can think of several improvements to this reaction:

The most important thing is to make sure that one uses exactly the right amount of chlorine. The production from TCCA+HCl is very cheap, but this reaction doesn't go to completion (that's why I used a bit more than the theoretical amount). Also, the commercially available material is only 92% TCCA. I will have to find out the amount of TCCA that produces exactly 0,1 mol of chlorine. I'll either have to collect the gas over water (via pneumatic bowl) or use it in a reaction where the product can be titrated.

Another important thing would be to use more concentrated solutions. Especially the NaOH can surely be used in a 50% solution (saturated). The urea can also be used as a more concentrated solution (the patent actually uses a slurry, but this would hinder the reaction since the mix is not homogenous, so I'll use a saturated solution instead).

When a high enough concentration of reactants can be reached, it might become economic to distill off the hydrazine as the hydrate instead of precipitating it.

Maybe I should, after optimization, publicate the entire process as "Mr. Anonymous 2"? :D

[Edited on 18-5-2005 by garage chemist]

Rosco Bodine - 18-5-2005 at 15:57

Outstanding that the chlorourea method is confirmed . It should be so much easier
to form the monochlorourea and not have to deal with the exotherm and possible overchlorination as with forming sodium hypochlorite . I haven't studied
the patent closely but it seems possible
to perhaps combine the methods , using
a lower concentration of bleach with some extra NaOH to provide an even more
concentrated reaction mixture . Whether
the reaction mechanism is basically the same and if this could work , I am not sure . This is simply a variation on the idea I have already mentioned about using the lower strength bleach in a different way , strengthening it by basifying it further and chlorinating it to a higher concentration . I have an idea that
sodium hypochlorite is actually formed in situ when the monochlorourea is added to the NaOH solution . Indeed urea forms
clathrates and the monochlorourea may be a clathrate , essentially nothing but chlorine soaked up in the urea which swells like a sponge entrapping it , and then releasing it upon dissolving in the aqueous NaOH .

The conditions of pH and temperature also provide opportunity for steering the final result to be semicarbazide or hydrazine . If hydrazine sulfate is the intended product from some of the more diluted reaction mixtures , it may be helpful to increase the amount of H2SO4
by ten or fifteen percent on a molar basis ,
as the solubility of the HS decreases rapidly in any excess of sulfuric acid .

chemoleo - 18-5-2005 at 16:03

Well done Garage Chemist!
I'd be really interested what yield you got, having done the Raschig process I know what a pain it is, and what large volumes have to be used!
I suppose one of the great advantages is the fact that small volumes are used from the beginning, avoiding this concentration step that is necessary with the Raschig process.

In other words, the hydrazine is already quite concentrated. I wonder if even smaller volumes for dissolution could be used. (edit:seems you mentioned this in an edit)

Oh, and why do you first use HCl to neutralise the NaOH, then H2SO4? Isn't it possible you have losses due to the more soluble hydrazine chloride? So why not using *only*H2SO4?


PS I hope you did a quick testrun before doing such large volumes - weren't you afraid of NCl3? ;)

Why also do you want to use exactly the right amount of chlorine? It might help to put up the reaction equation btw

Great work anyway!


[Edited on 19-5-2005 by chemoleo]

garage chemist - 18-5-2005 at 16:22

I used HCl for the first neutralisation because when only using H2SO4, one gets huge amounts of glauber's salt (Na2SO4*10H2O) along with the HS if one cools too deep. The solubility of NaCl doesn't change with temperature, which permits stronger cooling to get more HS.

If I use too much chlorine, dichlorourea would begin to form, or worse, NCl3 (I am certain that it forms from urea and chlorine, but only in acidic environment- the excess urea helps keeping the pH basic).

I added the first bit of chlorine very slowly and checked very carefully for any yellow color or oily droplets ( I know how it looks when NCl3 forms, having made it on purpose a long time ago- I also know the smell of it, which is VERY strong).

[Edited on 19-5-2005 by garage chemist]

Rosco Bodine - 18-5-2005 at 17:00

Just a couple of thoughts that may be useful , if the reaction mechanism is what I suspect it may be , and the monochlorourea / dichlorourea are unstable not well defined compounds chemically , but inclusion compounds , which are " sponges " for adsorbed chlorine . The Schestakow type reaction
is not adversely affected by excess NaOH ,
and the quantities I arrived at were for sufficiency and economy for the reaction using the sodium hypochlorite already formed in the usual way . But for the monochlorourea or polychlorinated urea
which may occur , it may be better to use a bit more total NaOH so that any excess chlorine is converted to hypochlorite . It is important to keep the ratio of urea to hypochlorite correct so there is no substantial excess of hypochlorite , and this should translate for the ratio of total chlorine to urea also . But it hurts nothing to have excess NaOH , it just requires more HCl for the intial neutralization . These ratios may vary some from what is best for the direct hypochlorite method may not be directly applicable to the monochlorourea method , but I suspect the ratios will be close enough that interchangeability is a valid starting point . Anyway , please do
try using the weaker bleach with NaOH ,
and then you need not worry about overchlorinating the urea , because you can have the unreacted urea present in an amount which will react with the hypochlorite already present in the bleach . This will avoid any danger for overchlorinating the urea , and at the same time will give a more concentrated reaction mixture .

When the foam is subsiding and you see
that happy bright orange color ......

Whew , I think it's gonna be alright
Yeah, the worst is over now .....

My reaction flask is shinin'
like a red rubber ball :D;):D;):D

All genius is seasoned with a sprinkle of madness ......

and for some of us , there is more of a sprinkle than for others :D

[Edited on 19-5-2005 by Rosco Bodine]

Attachment: redrubberball.mid (38kB)
This file has been downloaded 1505 times

Chlorine Simple Method

Rosco Bodine - 18-5-2005 at 18:53

Thanks to Organikum for this file mentioned in another thread .

Chlorine is evolved from the reaction between TCCA and ordinary salt ,
forming sodium cyanurate as the byproduct .

How best this reaction might be used in hydrazine synthesis may be a basis for some interesting experiments .
It may be that a slurry of salt , urea , and TCCA would form the monochlorourea .
And then the mixture added to basified bleach could perhaps form hydrazine
which may precipitate immediately as the
slightly soluble hydrazine cyanurate .
This could be then filtered out or perhaps converted in situ to hydrazine sulfate
by adding sulfuric acid to the heated
and stirred mixture . Actually the cyanurate can be freebased so it may be
pointless to convert to the sulfate , if
the cyanurate is more desirable for freebasing . This would depend upon the solubility of sodium cyanurate in alcohol . The possibilities associated with the usefulness of the cyanurate in azide synthesis would need to be considered ,
including its potential to be used directly .

After I attach this patent for the Chlorine from TCCA and salt , I will make a subsequent post with the file attached regarding the hydrazine cyanurate . Hopefully it is clear what I am describing as the reason these reactions are related , and how the reactions may possibly be combined in a synthesis of hydrazine .

[Edited on 19-5-2005 by Rosco Bodine]

Attachment: GB1401120 Chlorine from TCCA and salt.pdf (336kB)
This file has been downloaded 1723 times

Hydrazine Cyanurate

Rosco Bodine - 18-5-2005 at 19:21

GB1073292 Hydrazinium Cyanurate

This is another means of precipitating a slightly soluble hydrazine salt which may be freebased by NaOH .

Attachment: GB1073292 Hydrazinium Cyanurate.pdf (259kB)
This file has been downloaded 1656 times

garage chemist - 19-5-2005 at 03:11

The generation of Cl2 from TCCA and salt is very interesting, I haven't heard of such a reaction before. This will be one of my next experiments.
The precipitation of hydrazinium cyanurate from the mixture seemed like the ideal method to me first, but in the paper, it was also mentioned that monosodium cyanurate is insuluble in water. I'd be concerned about the purity of the precipitated HC, especially when an excess of NaOH was used.

From my used amounts of reactants, you can see that I used a bit over 0,3 mol urea, only 0,1 mol chlorine and 0,4 mol of NaOH. This provides an excess of urea and a further excess of NaOH and therefore absolutely no risk of urea overchlorination.

I don't think that the urea simply "adsorbs" the chlorine, because there was a strong chloramine (not chlorine!)- type smell over the chlorourea mixture.
The chlorine reacts with water to form HCl and HClO (an equilibrium, but it is strongly shifted to the right due to the basic environment). The HCl is immediately taken up by the urea to form urea hydrochloride, the same as with the generation of urea nitrate, just with HCl (and it stays in solution).
I'm sure that the HClO immediately reacts with urea to form N-chlorourea, because HClO is a very unstable compound (think of the rapid decomposion of acidified bleach!) and also very reactive.
You know the mechanism of the Hofmann degradation?

And you can be SURE that I'll never use any NaOCl products again. Not even as a supplementary chlorine donor in the chlorourea process. 100ml of my bleach only contain 2,8g NaOCl. You can't say that this would make any noticeable difference. It would just dilute the solution.

My next step will be the analysis of the chlorine- producing reaction and the determination of the required reactant amounts to make exactly 0,1 mol chlorine.
The TCCA/salt method will be used for this.

EDIT: The TCCA/salt method requires a vacuum. Therefore it isn't that useful.
HCl/TCCA will still be required.

[Edited on 19-5-2005 by garage chemist]

garage chemist - 19-5-2005 at 05:17

My HS weighs 6,6g.
That's an approx. 50% yield from chlorine.

Rosco Bodine - 19-5-2005 at 07:39

I see what you mean about the 2.8% bleach not offering any advantage because of the dilution . But the idea
of combining the reactions could be useful
if you have easy access to stronger bleach .

The fumes which come off the chlorourea reaction should be avoided , it being unknown what exactly may be the toxicity
or precise nature of the odor .

The reactions which may form hydrazine are a balancing act against undesired side reactions which may also occur if the proportions of reactants are not optimal for the hydrazine reaction to prevail .

Excess urea will react with the hydrazine formed , resulting in a carbonamide and other byproducts . Even some of the CO2
reacts with and decreases the yield of hydrazine . Excess NaOH tends to mitigate the formation of such byproducts .

Ordinary household bleach is about 5% NaOCL . And about the same amount of salt is present also IIRC . What I am trying to say is that instead of using plain distilled water to make your concentrated NaOH solution for the decomposition of the chlorourea , to use the bleach and
an extra molar equivalent of NaOH with regards to the NaOCl present , and then
the added amount of NaOH you would use for decomposition of the chlorourea .
Simultaneously , use an extra amount of urea equivalent to the molar amount of NaOCl , and let that excess portion of the total urea remain as unchlorinated excess urea present in the slurry of urea being chlorinated to monochlorourea . When the chlorourea and unreacted urea mixture is added to the basified hypochlorite , the reaction should proceed to produce hydrazine from two precursors simultaneously , converting the monochlorourea as in your reported experiment , and also converting the unchlorinated urea by reaction with the hypochlorite . The reaction would be a marrying of the Shestakow reaction with its related monochlorourea reaction , which may actually be the same reaction
only having different means of bringing chlorine into the reaction . You see I suspect that sodium hypochlorite is an intermediate in the reaction by either route , whether it is formed in advance ,
by using bleach and urea for the reaction ,
or whether it forms in situ when chlorourea reacts with NaOH . The chlorourea reaction for hydrazine should be more exothermic and more rapid because the heat of reaction forming the hypochlorite is added to the usual exotherm for the hydrazine reaction which occurs subsequently . This theory may be all hogwash , and I have not studied this nor done any experiments to
test the theory . It just occurred to me that it is worth examining as a possibility .

With regards to the TCCA and salt reaction needing vacuum , not necessarily so if there is something present or added gradually which will react with the free chlorine and shift the equlibrium to the right . Anything which reacts with the chlorine produced could emulate the effect of vacuum and cause the TCCA and salt reaction to proceed . Again I have not tested this theory to see if monochlorourea would form , but it is a worthwhile possibility for experimentation .

All of these ideas are speculations for methods of concentrating the reaction mixture producing hydrazine and simplifying the preparation of the precursors , and some or none of these ideas may prove to have any validity or be workable . These ideas just seem to present some new angles to explore in regards to experiments with hydrazine , where TCCA is the new reactant whose potential usefulness is yet to be determined . The truth is that it would take a series of experiments to actually determine what are the advantages if any for applying any of these untested ideas .

[Edited on 19-5-2005 by Rosco Bodine]

garage chemist - 19-5-2005 at 08:04

Do you think that 0,7g NaOCl (from the 20ml bleach used to dissolve the NaOH) with 20g urea have any measurable effect?

Using a stronger NaOH solution would raise the concentration of hydrazine in the resulting liquid much more than using bleach.

Rosco Bodine - 19-5-2005 at 08:16

Your message and my edit just passed each other in flight through cyberspace .

I understand exactly what you are saying and you are correct that with that low a concentration of bleach , there is likely no gain and probably a loss for trying to combine the methods .

Europeans need to riot in the streets demanding that their bleach be watered down no longer , that you are just not going to take it any more , paying top Euro for chlorox that has been " stepped on " by the middle man :D

But for those of us who can get 10% bleach at the grocery store , the possibility
of in effect doubling the concentration of the reaction mixture for the usual process
does have a certain appeal .

ordenblitz - 7-6-2005 at 17:47

I decided to try the chlorourea route to hydrazine sulfate today.
Generally I followed Garage Chemist’s method but made a few slight changes in temperature, timing and chlorine source.

In a 250ml beaker was added:
0.2 gm gelatin (Knox)
20.5 gm urea (Fisher a.c.s.)
45 ml water, R.O.
The beaker was warmed over a flame until contents dissolved then placed in an ice bath to chill.

A chlorine generator was cobbled together using a flask and addition funnel pretty much the same as Mr. Chemist’s set up. There in was placed:
12 gm of CaOCl (HTH containing 78% CaOCl & 22% unknown adulterant)
10 ml water, R.O.
30 ml HCl (hardware store brand 31%)

The chilled ( ~4c) urea/gelatin soln was placed in a 100 ml graduated cylinder and the chlorine was bubbled in through a glass pipette connected to generator using rubber tubing. The HCl was added slowly drop wise to CaOCl over a period of 15 min. The addition of the chlorine caused the temperature of the solution rise to about 30c. After the chlorine stopped coming over, the pipette was pulled out and the cylinder and contents allowed to rest. It was a lovely pale yellow color.

In a 1000 ml Erlenmeyer flask the following was added:
40 ml water R.O.
16 gm NaOH (Fisher NF/FCC)
2 in. octagon stir bar.
The flask was placed on a stirring hot plate and stirring begun. After the sodium hydroxide had dissolved, the temperature had only risen to ~55c due to the heat loss to the large flask.

The chlorourea was swiftly poured, in one portion, to the flask with the NaOH while stirring fast. The contents immediately foamed up to occupy at least 70% of the volume of the flask but there was no appreciable color change except for the lightening of color due to the foam. Stirring continued for 5 minutes further where upon the foam began to subside and the color turned slightly amber. Stirring was discontinued while the hotplate was turned up high and the flask covered with plastic wrap to keep out excess air. At this point the solution was no longer foaming at all and the temperature was probably 75c or so.

The flask was taken off the hot plate and left to rest, covered for 5 minutes then placed in a bowl of cool water. Ice was added then to further speed the cooling. After the contents cooled to about 5c it was put back on the stir plate and adjusted to a medium speed.
33 ml HCl (31.45%) was added in a slow stream with considerable frothing and white fumes, over the period of 1 to 2 minutes. Next, 21 ml H2SO4 (Fisher 95%) was added in a very slow stream in a few portions while trying to keep spattering to a minimum. The crystals began to form but not as rapidly as I would have expected but after a minute the crop was clearly good.

Once again the flask was removed from stirring and quickly cooled in an ice bath to 3 or 4c. The contents vacuum filtered and the supernant liquid used to wash the remaining crystals from the flask. A few ml of cold H2O was used as a wash then the vacuum was continued to remove as much excess moisture as could be gotten out.

The product was weighed at 12.83 grams while still slightly moist. I will weigh again after it is dry but expect the amount to be over 11 grams.

The whole process was straight forward, easy and I must say, fun to do. It took less than 1.5 hours to complete including set up and some clean up afterward. I plan on running a sample on the FTIR and comparing it to a sample of HS I made with reagent grade hydrazine hydrate and reagent grade H2SO4. I will post the results afterward.

chemoleo - 7-6-2005 at 18:50

Interesting. So you got fourier transform infrared spectoscopy available?
You know what you should do with this? Analyse the putative DPPP, the thread of which you surely must have noticed!

[Edited on 8-6-2005 by chemoleo]

ordenblitz - 13-6-2005 at 19:03

The final yield of HS from the chlorourea process was 9.62 grams. My guess was high as the crystals must have contained more water than I had guessed, but still a good quantity.

Here are a few pictures. The pan on the right contains high grade HS from a known source. The pan on the left is the material from chlorourea process.

CompareCry.JPG - 35kB

ordenblitz - 13-6-2005 at 19:10

The spectra in red is the material from the chlorourea process and the one in green is from my known source which is at least tech grade or better. Unfortunately my libraries did not contain a commercial spectra of HS so we all will have to compare by eye. The peaks under 3200 and 3400 are probably from residual H2SO4. I need to recrystalize and run it again but it looks like a good match from my experience.

HSspectra.JPG - 30kB

ordenblitz - 13-6-2005 at 19:13

Sorry about the multiple posts but I can't seem to figure out how to post more than one pic at a time with my old Mac.

Here are the starting materials for my HS from chlorourea experiments.

PrecHSTst.JPG - 36kB

ordenblitz - 13-6-2005 at 19:16

Where I got my known good hydrazine sulfate from.

KwnHSPre.JPG - 31kB

garage chemist - 14-6-2005 at 07:53

Good, work, ordenblitz!

But why was your yield so much higher than mine?
I suspect that you added more chlorine, since my chlorourea solution was only faintly yellow and your solution strong yellow.
Be careful though, NCl3 has exactly this intense yellow color...

Rosco Bodine - 14-6-2005 at 08:24

The modification of the chlorourea process for semicarbazide as the end product should also be interesting .

There has been mention of the possibility
of using semicarbazide as a precursor to azides , but I know of no references nor details concerning the reaction .

Earlier in the thread there was mention about the possible danger of nitrogen trichloride being a byproduct with the chlorourea process . It was mentioned in the pdf regarding the use of TCCA as an oxidant on page 2 , that it is possible for nitrogen trichloride to form in strongly acidic solutions of TCCA , so this may need to be considered when the source of chlorine is TCCA . Whether the formation of nitrogen trichloride is temperature dependant or not I do not know for certain . It would probably be a hazard avoided simply by conducting the chlorine generation at warmer temperatures at which any NCl3 formation or accumulation would not be favored , and evolution of Cl2 would be
assured . Curiously this is likely one of those reactions where heating is the essential guarantor of safety , and cooling would not be a good idea for the TCCA based " chlorine generator " :D

A parallel is likely true for the chlorourea formation . Keeping the temperature slightly warm , maybe 25 C for example ,
would be safer than say using an ice-salt
cooling for the chlorination of the urea .

[Edited on 14-6-2005 by Rosco Bodine]

hydrazine synthesis continued

hannibal - 16-6-2005 at 20:36

sorry, for starting this new thread......somehow i could not post my reply on the previous thread on "hydrazine synthesis" , apparently it has been closed.
in one of the previous post by NickF he proposes that use of calcium hypochlorite increases the total conversion of hydrazine.
I am quoting him
"I'd recommend using urea instead of ammonia for making hydrazine. It means you don't get a house full of ammonia, especially if you can't get very good ventillation!
Check out the Hive for a synth from urea. I tried it, using fertiliser grade urea, <5% NaOCl household bleach, gelatine, drain-cleaner NaOH granules and tap water. The yield was just over 15%, while this is not very good it is a very cheap method. The yield is improved greatly by using Ca(OCl)2 instead of NaOCl, because of the Ca(OH)2 present. Adding Ca(OH)2 to a synthesis using NaOCl also increases efficiency. Using distilled water helps by reducing the ammount of transition metal ions, and using pure urea would also be a great help. The text I followed reported a 75% yield based on urea!
To make azides, it's probably best to first make a solution of hydrazine nitrate, by mixing strong solutions of hydrazine sulphate and calcium nitrate, and then filtering off the CaSO4 ppte. Otherwise you might end up with a lot of ppted PbSO4/Ag2SO4 in your product, depending on what you are trying to make.
I wanted to make semicarbazide for NTO and azides, but never got a very good yield from any of the methods I tried. For making azides, it's MUCH easier to use hydrazine (and semicarbazide is not very much less toxic than hydrazine anyway). It would be nice to be able to make NTO, but because of the trouble of having to make the semicarbazide it's better to use another HE." ---- by NickF
If this is true then why use of sodium hypochlorite is recommended,although it is more expensive than calcium hypochlorite(because calcium hydroxide is cheaper than sodium hydroxide)......if we use calcium hypochlorite there is one more advantage that i see calcium carbonate formed in the reaction can be easily separated
from the hydrazine solution(because solubility of CaCO3 is less),this decreases the amount of sulphuric acid needed in the precipitation of hydrazine sulphate.
2Ca(OH)2+2NH2CONH2 +Ca(OCl)2---->2N2H4 + 2CaCO3 +CaCl2 +2H2O

Although the hydrazine sulphate that we get will be slightly more impure because of calcium sulphate precipitation in this case.
one more point why in this case a bivalent ion (Ca2+) does not catalyse the oxidation of hydrazine....if it does then why hydrazine yields are high.

Rosco Bodine - 17-6-2005 at 06:39

Quote:
Originally posted by hannibal
If this is true


* If * is a very big word in this entire matter . Suppose that the use of Ca hypochlorite is not superior , then what
good are the rest of the rational analyses ? No one has described the
details of a step by step method using
Ca hypochlorite with comparisons of the
yield with the sodium hypochlorite method , which is well proven and gives good yields . Ca hypochlorite is NOT cheaper than sodium hypochlorite , except
for those industrial scenarios where the
hypochlorite solutions are made on site for
immediate use . When carried to the point of making stable products for distribution and sale , the sodium hypochlorite solutions are cheaper to produce . Compare the OTC price of a gallon of 10% sodium hypochlorite for $1.50 , with an amount of granular HTH having the same chlorine content and you'll see very quickly the " dry chlorine "
product is several times more expensive .
Quote:
then why use of sodium hypochlorite is recommended,although it is more expensive than calcium hypochlorite(because calcium hydroxide is cheaper than sodium hydroxide)......if we use calcium hypochlorite there is one more advantage that i see calcium carbonate formed in the reaction can be easily separated
from the hydrazine solution(because solubility of CaCO3 is less),this decreases the amount of sulphuric acid needed in the precipitation of hydrazine sulphate.
2Ca(OH)2+2NH2CONH2 +Ca(OCl)2---->2N2H4 + 2CaCO3 +CaCl2 +2H2O

Although the hydrazine sulphate that we get will be slightly more impure because of calcium sulphate precipitation in this case.
one more point why in this case a bivalent ion (Ca2+) does not catalyse the oxidation of hydrazine....if it does then why hydrazine yields are high.


Anyone is welcome to test variants and chart the results of what improvements are
* proven * by alternate hypochlorites ,
mixtures of hypochlorites , and modified conditions for the reaction which * prove *
useful . But since it is known for certain that the urea process is sensitive to small variables , then there is no basis for any conclusions about what is better in the way of changes to the known process ,
before actually testing and retesting the modification to make certain it has validity . Knowing how sensitive the reaction is to small variables is exactly why I performed a half dozen experiments for confirmation , before having the boldness to declare that use
of hydrochloric acid for the initial neutralization , and then following that with sulfuric acid for securing precipitation of the sulfate , was a * superior * method . You see there is a huge *inertia* associated with textbook methods which have been in practice for a hundred years , and any maverick suggesting changes should be sure they aren't calling a press conference in Utah to report their success at cold fusion using a method nobody else on earth can reproduce :D

[Edited on 17-6-2005 by Rosco Bodine]

hannibal - 18-6-2005 at 00:43

after going through the entire discussion i still have some questions regarding synthesis of hydrazine from urea and sodium hypochlorite process ----
1.why are we adding excess NaOH in the sodium hypochlorite solution before adding urea into it and then take pain to again cool it to keep it stable.....a better way would be to add NaOH along with urea into hypochlorite.This will also decrease the amount of heat that we need to supply during the hydrazine formation.
2.why are we adding urea to hypochlorite and not hypochlorite to urea.
i think for somewhat larger batches adding of hypochlorite to urea would be better.....has someone tried it by adding hypochlorite to urea???????
3.we all are measuring the yield of hydrazine sulphate and not actual hydrazine formed (some amount remain dissolved in the solution), this method will give lower conversion and also conversions that cannot be compared as we all are dealing with different volumes of solutions.There are also some other complications like sodium sulphate precipitation along with hydrazine sulphate........i think there is a method for quantitative analysis of hydrazine by titrating it with iodine although i dont know the exact equations involved but know its pretty simple, the exact method for titration should definitely be in "hydrazine and its derivatives".Please post this method if anyone gets it

[Edited on 18-6-2005 by hannibal]

The_Davster - 18-6-2005 at 05:55

What does it matter for anything other than anylitical chemistry the 'real' yield of hydrazine formed? What matters to us and others in practical chemistry is the ammount of product that can be isolated from the reaction mix. Besides, using HS solubility we can figure out aproximatly how much is still in the solution.

@ hannibal

Rosco Bodine - 18-6-2005 at 07:10

There is nothing arbitrary about the methods which have been worked out for the synthesis of hydrazine , by methods which are practical for small scale production . The literature including the patents will give you the reactions which apply , and you can justify the proportions of reactants and their order of addition , and conditions of temperature
and concentration required , by experiments and comparisons of your
observations and results . An understanding of the chemistry and the properties of the materials used in the reaction , would set aside some of your questions before they are asked , such as why the base isn't mixed with urea , which would decompose the urea . But you can try it and see and also see whether that improves the production of hydrazine . Also reversing the order of addition you can try and see what happens . What we are trying to do is to pin down the details and refinements upon methods which are known to work well , not so much to explore every possible variation which imagination can provide , including those which seem contrary to theory . It requires enough time and work simply doing the experiments which follow from the logic which can be best defended . So there is little enthusiasm for experiments which look for the unexpected anomaly rarely found when a synthesis is deliberately run in a direction against theory , simply for satisfying curiosity as to what will be the result . Usually what happens is you either get none of the intended product at all , or a mixed product with all the possible byproducts
which may occur in a mixture which hasn't been deliberately steered to favor one specific target product from controlled reactions .

Hydrazine From Urea

MadHatter - 18-6-2005 at 17:36

Hannibal, the production of hydrazine from urea is more efficient, and less smelly, than
from ammonia. As far as your newer thread on this subject, The Bird has the final word !

hannibal - 20-6-2005 at 04:03

rosco i dont think i have written anything illogical ..... regarding mixing of urea and the base issue let me make myself clear ..... i am not mixing sodium hydroxide in urea and then adding it to hypochlorite instead i am adding urea and sodium hydroxide separately but simultaneously in the hypochlorite.pls tell me if i have made myself clear this time.
anyways i am planning to do some experiment somewhere in the mid of next month to see the results myself...any suggestions???????

Rosco Bodine - 20-6-2005 at 06:58

hannibal ,

The suggestion I would emphasize is that the first time you make hydrazine , in order to see exactly what sort of reaction dynamic which occurs pretty rapidly and in a narrow range of conditions , first duplicate one of the methods that have already been worked out and proven . You will then have an insight into some factors which make these labscale batch processes necessary to be done the way they are done , to simplify things and to get a decent yield by a lab method that is
convenient . Many of the things which could be done in an industrial process are
not easily adapted to a batch process or
produce so little increase in yield as to not be worth the trouble . On an industrial scale with continuous separate feeds of separate precursors , where controlled flow rates and reaction temperatures , residence times and other
parameters can be controlled , it is easy enough to explore many variations which can tweak the reaction and get a few per cent increase yields . It is much more difficult to implement manageable changes to a vat process which is already
optimized for a lab scale . It is even more difficult to do that for a reaction that you don't appreciate is more complicated than
just a linear reaction like A+B+C=D . Hydrazine is actually a competing reaction with a lot of other possible reactions , and if you don't hit the target conditions just right , then you won't get a good yield of hydrazine . Concentration of reactants is a factor which makes the solubility of the precursors a very pertinent concern , so the low solubility of
CaOH does not fit the right thinking that you need a very soluble base for a concentrated reaction mixture , aside from the extra steps introduced by having to separate calcium byproduct precipitates . When you have experience with the reaction that works , then you will understand that some of the things you suggest actually aren't logical . Do some experiments to prove that if you need the proof . When you get better yields than 65% of dried crystals of pure hydrazine sulfate , then describe the details of what works better than what we already know and there won't be any speculation about what is an improvement
* after * the details are known . Debating the details in advance of the finding is premature .
I can make kilos of hydrazine if I want it , while you are trying to analyze why your failed experiments didn't work . There's the point . As for using the heat of dilution of the NaOH to be an enhancement of the exotherm for the production of the hydrazine , well that would make sense if the reaction needed the boost . Simply varying the initiation temperature upwards a few degrees will give you all the boost you can handle , especially for larger batch sizes . So you will never need the added exotherm from such a strategy . Also there would be the difficulty introduced of getting a consistent mixture , before the foaming complicates any further addition , which sometimes occurs almost instantly , especially at warmer temperatures for the start of the reaction .

[Edited on 20-6-2005 by Rosco Bodine]

hannibal - 20-6-2005 at 20:42

rosco let us talk ablout these things after i complete the experiments and believe me my aim would not be to prove u wrong but to give an improved synthesis....can anyone tell me about the reaction of hydrazine with iodine in sodium bicarbonate...may be madhatter u can tell from "hydrazine and its derivatives"

Rosco Bodine - 21-6-2005 at 09:34

hannibal ,

The biggest potential improvement I see possible in terms of increasing the yield of hydrazine per unit volume of reaction mixture , is in combining the chlorourea method with the urea-hypochlorite method as I mentioned earlier in the thread . Concurrent isolation of the hydrazine as a hydrazone intermediate might be another possible enhancement ,
but whether these things are cost effective in terms of the added steps and reagents , over the straightforward single methods is unlikely . It is probably better just to increase the batch size . That's another real consideration for reactions that are producing yields at 65% , where you may be able to improve that to 75% or 80% , but the tradeoff is that the added work or cost of added reagents , or both combined , makes it easier and more economical just to increase the batch size for the lower yielding method , and be content that
65% is " good enough " , if you follow the rationale .

[Edited on 21-6-2005 by Rosco Bodine]

Hydrazine And Its Derivatives - 1st Edition

MadHatter - 21-6-2005 at 20:26

Bear in mind that my book is 1st Edition and published in its 1st year - 1984. The 2nd Edition,
which I can't find at reasonable price for the moment, is twice as long - even in the synthesis
section.

Hannibal, there is no mention of iodine or sodium bicarbonate being used in the urea process -
at least not in the 1st Edition. I don't know about the 2nd Edition because I haven't
seen it - yet ! This edition states the maximum efficiency of the urea process is 64%. Maybe there's
been improvements discussed in the 2nd Edition, so Roscoe it sounds like your right on track.

Hannibal, don't be discouraged. Do some testing, and if you can improve the yields, we'll all
be looking forward to it !

BTW, I haven't been able to do much scanning lately. I'm setting up an FTP server with
200 GB capacity and I'm still working on the configuration. The 1st Edition will be the 1st book
I put there.

Analysis

MadHatter - 25-6-2005 at 19:17

Hannibal, sorry I misread your post about iodine and bicarbonate. I belive this is the information
from page 415 of Hydrazine And Its Derivatives, 1st Edition, 1984, that you're referring to:


" Hydrazine can be titrated directly with standard iodine solution if a
buffer is added to keep the pH between 7.0 and 7.4. The last drops of iodine
must be added slowly, and several seconds must be allowed for complete reac-
tion. At a pH lower than 7, the reaction is quantitative, but very slow. If
the pH increases above 7.5, hypoiodite formation may cause low results,
or hydrazine may begin to autoxidize. "

Bicarbonates, acetates, and tartrates are used as the buffers. According to the book, you can't
use too much buffer or the iodine won't be taken up - particularly true for sodium bicarbonate.
Also the last drops of iodine solution must be added very slowly. If the yellow color persists for
more than 2 minutes, the endpoint has been reached.


P.S. Still trying to get my FTP up and running. Still can't access from a remote computer.
More tweaking of the network and/or server settings necessary.

[Edited on 26-6-2005 by MadHatter]

a little help...?

mandarine - 19-7-2005 at 17:47

just interested in ending up with hydrazine hydrate here, not a question about converting it to the sulfate.

so, I've been at this reaction for a couple weeks now. all OTC, sodium hypochlorite (30ml of 5%), 10% janitorial strength aqueous ammonia (130ml of 10%), some gelatin (1.5 grams), and heating it to 1/3rd volume.

first of all, I'd like to clarify what on earth you end up with as soon as it has evaporated down to 1/3rd? I know there should be plenty of Chloramine, a little hydrazine hydrate, leftover gelatin, but that should be it? the solution is very much yellow, and because Chloramine is supposed to be yellow, I assume it's the most abundant result of the reaction.

now, I'm only interested in knowing that I can get this reaction right. I don't care about getting tons of hydrazine, I simply want to end up with a little relatively pure hydrazine. I'd even be happy with only a couple milliliters, frankly. in fact, at this point, just knowing FOR SURE that it's hydrazine and only having a milliliter would be REAL SWELL :mad: :D

what I've been trying from this point is to add 1 volume of xylene to the resultant liquid after it's cooled down to room temp. unfortunately I can't find any info on Chloramine's solubility in xylene, but I know that hydrazine is PARTIALLY soluble in it. good enough for me. so I shake it in xylene, and then siphon the xylene layer into a new flask and add a couple milliliters of deionized water and shake again. the idea, of course, being that there ought to be SOME hydrazine in that xylene, and hydrazine is far more attracted to water than xylene, so you'd think a few milliliters of water would combine with hydrazine and become a relatively high concentration of hydrazine hydrate?

the water is siphoned out and takes on a pH of 11/11.5 (?!) (which is light brown on regular pH paper). WTF? isn't hydrazine, in its ANHYDROUS FORM, only supposed to get up to 10 (green on pH paper)? with lower and lower pH as is becomes more dilute with water? obviously something's going wrong somewhere, can anyone help?

now, trying to reason through this, most of the ammonia and hypochlorite would have finished reacting after the reaction mixture hit 60C, and they surely wouldn’t be present anymore after it hits 100C (I usually heat up around 110C). so, all that we can be left with in the yellow mixture is what I listed above, and none of those things has a pH of 11 (chloramine's pH is 9)... I'm really confused here.

how stable is hydrazine to oxygen? does it only last for a couple hours when exposed to air or something, then decomposing into ammonia/etc? I know UV breaks it down too...

I've tried the same xylene extraction process with the liquid boiled down to 1/6th and 1/2. the final water extract still has a mysteriously high pH.

does anyone have good ideas for testing the hydrazine?

as I know it's supposed to react a little with copper, I dripped a single drop of my supposed hydrazine hydrate onto two pennies, one modern (only lightly coated with copper, 2.5%, the rest being zinc metal) and one from the 1950's (which is 95% copper and 5% zinc).

after sitting with a drop of the supposed hydrazine overnight, each penny turned very clearly green at the place where the liquid was, with the liquid of course having evaporated by then.

as those results mean nothing to me, I also tried dropping a single droplet of the supposed hydrazine on an open flame :P

nothing happened.

am I somehow just ending up with ammonia here or what? and if so, why?

*sigh* can anyone interpret these results? anyone have better tests for hydrazine (besides showering yourself in it and seeing if you get cancer)?

BTW, after turning the heat back on after it has foamed to evaporate down to 1/3rd, I’m always OUT OF THE HOUSE. for safety, all I do is perform the reaction next to a big window (there’s no people and hardly any wildlife around here) with a fan blowing directly on the mouth of the reacting flask and out the window. I also get a little dizzy from holding my breath, plugging my nose, and pressing goggles against my face whenever I have to be near the flask (which is as rarely as possible). am I being paranoid here?

and what about Chloramine? I've searched, but can't find any info on its solubility with xylene. it's obviously very soluble in water, as regular chlorine is being phased out in favor of Chloramine to treat public water. does anyone have a link to the chemical structure of Chloramine? Wikipedia usually is very good about putting the structure of chemical compounds with their chemical pages, but not for chloramine.

maybe someone has a suggestion for a fool proof method to extract hydrazine from this yellow liquid? and speaking of yellow liquid, where can one go to get OTC urea? I assume fasting for a day, drinking lots of water and peeing onto a hypochlorite isn't an advised method of getting hydrazine ;)

and here's another question: would adding some gypsum salts, epsom salts, or other drying agents along with the gelatin be a good idea for ending up with less water and more hydrazine? probably not?

and here’s on FINAL question (sorry for all the questions!). I’ve heard that only anhydrous hydrazine can react with an amine to produce an azide (when considering only the hydrazine family anyway). is this true? can hydrazine hydrate work to produce an azide from an amine too? what about if it's very dilute?

P.S. all non-OTC suggestions, while appearing very interesting to me, will not help anything ;)

[Edited on 20-7-2005 by mandarine]

[Edited on 20-7-2005 by mandarine]

[Edited on 20-7-2005 by mandarine]

neutrino - 19-7-2005 at 18:26

Hydrazine is very reactive toward oxygen. In fact, one of its main industrial uses is as an oxygen scavenger in boilers because it reacts so well.

I'm not sure about amines, but nitrites are the standard reagents we use to make azide around here.

BromicAcid - 19-7-2005 at 18:26

Wow.... lots of questions.... I'll answer what I can. First off, you could just attempt to distill hydrazine hydrate from your mixture for making hydrazine in the first place paying careful attention to the fractions and collecting at the boiling point of the azeotrope, I have read of this being done. Also note that hydrazine is recomended to be distilled under hydrogen in a copper or silver apparatus, it will react with glass to the best of my knowledge (though my distillation attempts of hydrazine hydrate show no attack on the glass) also note that ground glass with hydrazine on it can explode, the high surface area working catalytically with oxygen in the air. Hydrazine will react with the oxygen in the air anyway under its own power with little provolking.

Now, you are trying to see if you have hydrazine by checking the pH of the solution you get from extracting with xylene and adding water, hydrazine is actually classified as a somewhat strong base, anhydrous hydrazine might have a lower pH then the hydrazine hydrate, same methodology that ~75% sulfuric acid has more hydronium ions then 100% sulfuric acid. So the pH method is not a good test for hydrazine.

As for tests for hydrazine there are little pads of paper that you can put water on containing hydrazine which will detect in the ppm but are otherwise useless, you could test it's reducing power using the silver mirror test or something similar but there are other qualitative tests for hydrazine, a few are in this thread if you read it carefully.

ordenblitz - 19-7-2005 at 19:10

Mandarine,

IMHO the best way, is to proceed as directed and convert your product to hydrazine sulfate. You will then be able to determine easily how much hydrazine you actually made. If you go back on the thread you will find directions on freebasing with NaOH back to hydrazine. Many in the explosive industry purchase their hydrazine as the sulfate and recover as needed. It's the safest and best way to store NH2NH2.

[Edited on 21-7-2005 by ordenblitz]

Hydrazine

nikcorbet - 9-9-2005 at 05:44

What MP of hydrazine sulfate you get by Urea process . can any one tell actually it should be 254 Deg C but its not that after crystln twice :mad:

[Edited on 9-9-2005 by nikcorbet]

redneck - 1-12-2005 at 02:54

Has anybody any expierience with this synthesis?


Working up synthesis solutions obtained in production of hydrazine , United States Patent 4005179


Abstract: In the production of hydrazine wherein aqueous ammonia is oxidized in the presence of a ketone to form an aqueous solution containing at least one of a hydrazone and a ketazine along with ammonia, the hydrazone and ketazine are concentrated and the hydrazone and ketazine are subsequently hydrolyzed to hydrazine and ketone, the improvement which comprises effecting the concentration of the hydrazone and ketazine by extracting the aqueous solution with a substantially water-immiscible organic solvent whereby the hydrazone and ketazine preferentially enter the water-immiscible solvent, and separating the water-immiscible solvent extract from the aqueous solution. The organic solvent is preferably a higher alcohol, a chlorinated hydrocarbon, benzene or a substitution product thereof. The organic solvent extract, in one or more stages, and optionally with addition of ketone, can be re-extracted with water, hydrolyzed with aqueous acid or subjected to distillation to separate the hydrazine values from the organic solvent.

see:
http://freepatentsonline.com/4005179.html

Search

chloric1 - 1-12-2005 at 09:02

I am pretty sure this method was already dscussed. Try searching the forum and I am sure you will find something. I remember writing about mekazine.

I cannot remember the patent number but the ketazines can be produced from ketones, ammonia and hypochlorites. THis should be easier for the home chemist. You just mix 2 moles of ammonia fir every mole of hypochlorite to form your chloramine(preferably at 0C) and add this solution to a well stirred mixture of MEK and concentrated ammonia. Gelatine is optional. Ultimately you would use about 10 to 15: 1 ratio for ammonia to hypochlorite. The ammonia to ketone should be more than a 4:1 ratio.

I will let you know if and when I find the patent.

P.S. the ketazines hydrolysis in water is under 10 atmospheres:(. When strong acids are used, it happens at standard pressure and only a little heat is required.

jester - 12-12-2005 at 21:17

has somebody posted a link for hydrazine and its derivatives yet

stricnine - 23-12-2005 at 04:47

Hi there!

Hydrazine was used by the German Luftwaffe to power the Messerschmitt Me 163, an interceptor (II WW!) powered by a hydrazine-hydrogen peroxide rocket (clearly based upon the fact that it reacts readily with oxygen). The thing reached over 1000Km/h in 1944/45! Perhaps this could start a new thread in the energetic reactions forum!

Cheers

Stricnine

verode - 8-8-2006 at 08:56

also urea and hypobromite or hypochlorite you get (hoffman) the hydrazine

Theoretic - 9-12-2006 at 07:34

Me 163 used hydrazine hydrate + methanol as fuel.

Sauron - 22-12-2006 at 21:41

The Org.Syn. procedures for hydrazine, hydrazine hydrate, hydrazine suphate, and neat anhydrous hydrazine have stodd the test of time and are available freely online.

Hydrazine in any form is a nasty beast so I hope you have SCBA or at least a good hood. Where I am one can buy the sulphate and hydrate freely and in bulk. Generally these are regarded as proven human carcinogens (which are few and far between) just in case you were not aware.

Rosco Bodine - 26-12-2006 at 15:38

Quote:
Originally posted by Sauron
The Org.Syn. procedures for hydrazine, hydrazine hydrate, hydrazine suphate, and neat anhydrous hydrazine have stodd the test of time and are available freely online.


The method described in Wiley for hydrazine sulfate is in every way inferior and primitive compared to the two alternative methods described in this forum for synthesis from urea .

Sauron - 26-12-2006 at 16:10

I believe the Hoffman is best done with hypohalite freshly prepared in situ because off the shelf solutions always contain appreciate amounts of chlorate ion.

I tried this rxn a couple of times on pthalimide (trying to obtain anthanilic acid) but all I got was a gray-black gunky mess.

If memory serves doesn't this rxn proceed through the isocyanate?



[Edited on 27-12-2006 by Sauron]

Rosco Bodine - 26-12-2006 at 16:49

Read this whole thread . There are other related threads
and plenty of patent references cited .

I personally developed a lab scale optimized adaptation of the Merck patent process ( GB392845 ) from many experiments , which have been confirmed by others .

http://sciencemadness.org/talk/viewthread.php?tid=757#pid656...

Going beyond that , the chlorourea method has been described earlier in this thread . The hydrazine sulfate
from both methods has been used to produce sodium azide , and also hydrazoic acid IIRC , via reaction with H2O2 ,

http://sciencemadness.org/talk/viewthread.php?tid=1987&p...

as well as having structure confirmed instrumentally by FTIR spectroscopy by ordenblitz here in this thread .

We probably should accumulate all the hydrazine / azide related experiments into one ScienceMadness " correlative volume " and put it in the Forum Library as well as send it to Wiley as an addendum / update to their antiquated
and obsolete method .

[Edited on 27-12-2006 by Rosco Bodine]

S.C. Wack - 26-12-2006 at 21:36

Sauron seems to have deleted where he doubted making hydrazine from urea and hypochlorite, which makes my looking up DE164755, US1959503 (the above GB patent), DE729105, and DE735321 seem pointless.

Sauron - 26-12-2006 at 22:05

I will happily look at those patents.

I deleted a hasty and excessive skepticism about a topic off my patch.

Here, hydrazine hydrate and sulphate are reradily purchaseable in both lab and bulk qty. No one seems overly worried about its toxicity or about damnfools making astrolite. My main supply jouse has 20 L pails of the hydrate at 50% if I recall. I doubt it is expensive. I rarely have a use for it as I am not making bootleg Viagra and its former prominence in peptide work is passe.

Rosco Bodine - 27-12-2006 at 00:53

Hydrazine and a few other " essential reagents " have been topics of great interest for the experimenters who
do not wish to purchase such chemicals for whatever reasons , or who simply prefer to make their own reagents from the most basic precursors . Sometimes this is an academic exercise and revisiting of the work of early researchers experimentation , out of curiosity about
what may have been missed or gone unreported or unrecognized in significance , and also looking for what is very good knowledge but has become "lost art" in the present texts .

When you want to learn about grandpas practical
synthetic methods ....you have to talk to grandpa , or
read his journals and patents and try to glean from those what may still be very useful information today .

It never hurts either for the younger folks to understand that some *very* smart men lived long long ago . And those ancients figured out a few things which shouldn't need to be rediscovered in historical obscurity . Yet older work often seems to be rediscovered and reinvented and revalued in some new way when the technology and materials of the present is applicable in some beneficial way to methods and knowledge of the past . It can be
properly humbling to hear an appropriate "I told you so"
from pages written fifty or a hundred years or more ago :D ....written in answer to a question it was known
in advance that someone was going to sooner or later be asking and find already answered at least in part :D

This is what puts the " re " in re-search you know ,
often travelling a path travelled before , but having a
different pair of eyes observing the scenery and
seeing if more can be seen by this traveller than
was recognized by the one(s) who went before ,
and perhaps to add another cumulative list of
footnotes to mark the point of what more is known now ,
than what was the assembled knowledge before .

Sauron - 27-12-2006 at 01:43

"If we see far it is merely because we are standing on the shoulders of giants."

A scientist paying homage to those who went before. I forget who it was who said it. It's an immortal line.

And I hate to admit it but I am long since old enough to be a grandpa.

[Edited on 27-12-2006 by Sauron]

Rosco Bodine - 27-12-2006 at 02:22

Greetings to a fellow dinosaur :D

For I am a grandpa too .....

Just wish I knew when I was young what I know now ,
for life would have been much more fun and profitable
and at a time when I was still young and healthy enough to enjoy it .

By the time you are old enough to know the best of what life is all about .....the best of life is already behind you ,
and this is a cruel irony .

Sauron - 27-12-2006 at 06:11

Ain't it the god-awful truth? Like General Sternwood in "The Big Sleep" I'm a very dull survival of a very gaudy life. Now reduced to enjoying my vices only by proxy.

unionised - 1-2-2007 at 12:36

According to some source I read this "If we see far it is merely because we are standing on the shoulders of giants." was a dig by Newton at his contemporary Hooke who was short. but then again, maybe not.
http://en.wikipedia.org/wiki/Robert_Hooke

hydrazine precautions.

joe69cool - 2-3-2007 at 19:50

I recently read that anhydrous hydrazine can be prepared from hydrazine sulfate by treatment with caustic methanol. I'm curious, would it be safe to distill off the methanol or would that lead to an ugly explosion (I think it would)? I realise its not something to be done in the home lab, but I'm curious never the less. Thanks, and pardon my ignorance.

[Edited on 3-3-2007 by joe69cool]

BromicAcid - 2-3-2007 at 19:59

Not really anhydrous...

N<sub>2</sub>H<sub>4</sub>* H<sub>2</sub>SO<sub>4</sub> + 2NaOH<sub>methanol</sub> ---> N<sub>2</sub>H<sub>4</sub> + Na<sub>2</sub>SO<sub>4</sub> + <b>2H<sub>2</sub>O</b>

I have done this prep before using potassium hydroxide dissolved in methanol. It goes well. But after filtering off the potassium sulfate and distilling off the methanol you are still left with hydrated hydrazine in the end.

[Edited on 3/3/2007 by BromicAcid]

Furch - 11-3-2007 at 20:10

Try this and you'll rid the hydrated hydrazine:

2 MeONa + N2H4*H2SO4 -methanol-> N2H4 + 2 MeOH + Na2SO4

Sodium might be a little too expensive for the cause, though...

DeAdFX - 13-4-2007 at 18:43

Is it possible to use Urea in place of ammonia in the keatazine process? Something tells me that Urea and acetone aren't going to react to form keatazine in any useful amount.

I was thinking that I could mix urea, acetone, and water and reflux for a few hours. The water would be for the hydrolsis of urea --> 2NH3 + CO2.

not_important - 13-4-2007 at 19:33

Urea doesn't hydrolise that quickly with catalysts. You need to get above 60 C in plain water to see much happening, and it still is fairly slow. Given the boiling point of acetone the rate of hydrolysis is going to be real slow.

garage chemist - 14-4-2007 at 12:19

I have once again started experimentation concerning the preparation of hydrazine.
I have chosen to further develop the chlorourea process instead of any other processes, which use either materials with very limited shelf-life and therefore unknown exact concentration (NaOCl over 5%) or have not been found to work on a small scale (ketazine process).

I have found that when chlorinating urea in aqueous solution, the solution turns green and slowly gives off colorless gas when 1 mol chlorine is used for 1,5 mol urea, and foams like beer when poured or shaken- clearly signs of decomposition. It seems necessary to use not more than 1 mol chlorine per 2 mol urea, as the solution does not turn green until shortly before all chlorine has been passed in (with 1mol Cl2/1,5mol urea).
I used a 50% aqueous solution of urea for the chlorination.

The chlorourea solution has been introduced into 50% NaOH solution with cooling and very vigorous stirring. A dense white fog is produced, and some gas.

I have done two batches, each with 0,5mol Cl2/0,75mol urea and one using 1,5mol NaOH and one using 1,125mol NaOH, in order to find out if using less NaOH affects yield a lot, since saving NaOH also saves a lot of acid in the neutralization.

The alkaline solution obtained by combining the chlorourea solution from 50% urea in water and 50% NaOH solution was slowly heated to boiling until every green color had vanished.

The resulting solution was so concentrated that a substantial amount of NaCl crystallized out, more of it on cooling (The solubility of NaCl in this mixed alkaline solution seems to be temperature dependent, unlike solutions of NaCl in water- or the precipitate contains another salt, like sodium carbonate, which is another byproduct).

This would make it possible to directly distill a dilute hydrazine hydrate solution (its only ca. 250ml of solution for a 0,5 mol Cl2 batch, with theoretical yield of 65g hydrazine sulfate!) without acidification. This would be an extremely favorable way of isolating the hydrazine, as neutralization of the distillate with H2SO4 would provide the stable and easily stored sulfate in quantitative yield, as opposed to direct addition of H2SO4 to the previously neutralized (with HCl) reaction mix, which leaves some of the product in solution due to its considerable solubility even at 0°C.
I have not tried this direct distillation of hydrazine hydrate from the reaction. If the two batches give encouraging yields of the sulfate, this will be my next experiment.

[Edited on 14-4-2007 by garage chemist]

garage chemist - 19-5-2007 at 14:00

The yields of the chlorourea batches were terrible, totally useless.
First one gave ca. 15g, second one maybe 5g. I abandoned the chlorourea method for me, primarily because the chlorourea is unstable and a huge excess of urea over chlorine has to be used, among other reasons.

-------------------------------

Today I have prepared the first batch of hydrazine sulfate using 70% Calcium hypochlorite (HTH pool chlorine) as the hypochlorite source.

I used 51,1g HTH, which I reacted with 37,8g Na2CO3, both as solutions, in the way I described in this thread:
http://www.sciencemadness.org/talk/viewthread.php?tid=8504

190ml of 15% NaOCl were obtained, corresponding to 0,45mol of NaOCl, a 90% yield in regard to HTH.

The rest of the synthesis was analogous to Mr. Anonymous' method for hydrazine sulfate, with some changes implemented by me.

I noticed that the solid NaOH dissolved extremely slowly in the 15% NaOCl when it was cold, and only started to dissolve with any noticeable speed as the temperature went above 10°C. But then the temperature did not rise above 20°C after the portion of NaOH had dissolved, so there was no problem.
I added the NaOH in two portions of 22,5g each (with cooling between additions to 8°C), for a total of 45g NaOH with the 0,45 mol NaOCl.

The Improvement:
The 15% NaOCl was not chilled in the freezer, the fridge with 8°C did perfectly fine. With the NaOH added in two portions, the temperature will not go above 20°C.
The same goes for the original HS synthesis when 10% NaOCl pool chlorinator is used!
If the NaOCl is chilled in the freezer instead of the fridge, the
NaOCl will simply warm up by itself to 10°C before the NaOH starts to dissolve. So not cooling below 0°C makes the process both easier and faster.
The NaOH stabilizes the NaOCl, so that 20°C can be tolerated for a short period of time if the NaOCl is chilled immediately after again.

After all NaOH had been dissolved, the NaOCl was chilled in the freezer to below 0°C, urea (34,1g) and gelatin (0,56g) were dissolved together in 40ml warm water and the instructions of Mr. Anonymous were followed until the end.
The foam rose to 1000ml from the ca. 250ml of solution.
The reaction became extremely hot by itself after the foam was produced. It subsided again quickly after.

Slowly heated to boiling until colorless, cooled, neutralized with 120ml 31% HCl and precipitated with 34ml conc H2SO4 (previously diluted with its own volume of water).

The batch is currently cooling. A lot of HS was already precipitating after the H2SO4 had been added, despite the solution still being hot. I heated until all had been dissolved, and then slowly left it cool in order to produce larger crystals. This works extremely well, as I can already see.

[Edited on 19-5-2007 by garage chemist]

some reference tables

Rosco Bodine - 20-5-2007 at 00:05

Here is a handy reference table for density and concentration of bleach along with some other general information .

Another idea sort of came to me regarding a possible usefulness of TCCA , which might produce calcium cyanurate
as an insoluble byproduct , which might later be useful for
a pyrolytic decomposition to calcium cyanamide . I haven't really written anything down as an equation to see if it
looks like it would work in a scheme for sodium hypochlorite .
There are perhaps a couple of different ways to go about it ,
so that you end up with a salt and sodium hypochlorite (bleach solution) as the supernatant , with a precipitate of
calcium cyanurate as a useful byproduct .

I'm thinking maybe calcium chloride together with sodium hydroxide and TCCA in some sequence or stirred and digested together for a couple of hours . Or perhaps
hydrated lime and sodium carbonate or bicarbonate plus TCCA might work , with a gradual titration with HCl as it proceeds . I'm not really sure what are the chances of
something like this working , but it would provide an interesting synthetic route eventually to tetrazoles and friends . So it may be a brainstorm or a brainfart , depending
on whether it would work out ....but I couldn't help but grin a little when the idea occurred to me , as it would be pretty slick if it did work .

[Edited on 20-5-2007 by Rosco Bodine]

Attachment: bleach.pdf (80kB)
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garage chemist - 20-5-2007 at 07:22

It worked awesome, I got 34,45g hydrazine sulfate from 51,1g 70% HTH on the first try!

Thats a 59% yield based on the sodium hypochlorite obtained as the intermediate. And that with incorrect stochiometry, as I now found out (I used too much urea and too much NaOH), and rather small batch size.
The yield may become even better when using a larger batch size and correct stochiometry!

This method of generating a 15% NaOCl, low-NaCl solution from calcium hypochlorite and Na2CO3 is clearly a very good method for running the hydrazine-producing reaction in as concentrated a solution as possible.

Now it is no longer necessary to use liquid pool chlorinator as the raw material. The use of calcium hypochlorite as raw material has been proven to be an excellent substitute, with almost as good yields.
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